Bonding

What is a Bond?• A force that holds atoms together.• Why?• We will look at it in terms of energy.• Bond energy- the energy required to

break a bond.• Why are compounds formed?• Because it gives the system the lowest

energy.

Ionic Bonding

• An atom with a low ionization energy reacts with an atom with high electron affinity.

• A metal and a non metal

• The electron moves.

• Opposite charges hold the atoms together.

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What about covalent compounds?

• The electrons in each atom are attracted to the nucleus of the other.

• The electrons repel each other,

• The nuclei repel each other.

• The reach a distance with the lowest possible energy.

• The distance between is the bond length.

0

En

erg

y

Internuclear Distance

0

En

erg

y


Bond Length

0

En

erg

y


Bond Energy

Covalent Bonding

• Electrons are shared by atoms.

• polar covalent bonds: The electrons are not shared evenly.

• One end is slightly positive, the other negative.

• Indicated using small delta

Wait, you already know about polarity.

H - F+ -

H - F+ -

H - F

+-H - F+

-

H - F

+-

H - F +-

H - F+-

H - F

+-

H - F

+-

H - F+ -

H - F

+-H - F+

-

H - F

+-

H - F +-

H - F+-

H - F

+-

H - F

+-

+-

Electronegativity

• The ability of an electron to attract shared electrons to itself.

Electronegativity• Tends to increase left to right.

• Decreases as you go down a group.

• Most noble gases are inert, no electronegativity.

• Difference in electronegativity between atoms tells us how polar the bond is.

Electronegativity difference

Bond Type

Zero

Intermediate

Large

Covalent

Polar Covalent

Ionic

Co

valent C

haracter

decreases

Ion

ic Ch

aracter increases

Dipole Moments• A molecule with a center of negative

charge and a center of positive charge is dipolar (two poles), or has a dipole moment.

• Center of charge doesn’t have to be on an atom.

• Will line up in the presence of an electric field.

H - F+ -

H - F

+-H - F+

-

H - F

+-

H - F +-

H - F+-

H - F

+-

H - F

+-

+-

How It is drawn

H - F+ -

Which Molecules Have Dipoles?

• Any two atom molecule with a polar bond.

• With three or more atoms there are two considerations.

1) There must be a polar bond.

2) Geometry can’t cancel it out.

Geometry and polarity• Three shapes will cancel them out.• Linear

Geometry and polarity

• Three shapes will cancel them out.

• Planar triangles

120º

Geometry and polarity• Three shapes will cancel them out.• Tetrahedral


• Others don’t cancel

• Bent


• Others don’t cancel

• Trigonal Pyramidal

Polar molecules• Must have polar bonds

• Must not be symmetrical

• Symmetrical shapes include• Linear• Trigonal planar• Tetrahedral• Trigonal bipyrimidal• Octahedral• Square planar

Ions• Atoms tend to react to form noble gas

configuration.

• Metals lose electrons to form cations

• Nonmetals can share electrons in covalent bonds. • When two non-metals react.(more later)

• Or they can gain electrons to form anions.

Ionic Compounds• We mean the solid crystal.

• Ions align themselves to maximize attractions between opposite charges, and to minimize repulsion between like ions.

• Can stabilize ions that would be unstable as a gas.

• React to achieve noble gas configuration (to form an octet).

Size of ions

• Ion size increases down a group.

• Cations are smaller than the atoms they came from.

• Anions are larger

Periodic Trends• Across the period nuclear charge

increases so they get smaller.• Energy level changes between anions

and cations.

Li+1

Be+2

B+3

C+4

N-3O-2 F-1

Size of Isoelectronic ions

• Positive ions have more protons so they are smaller.

Al+3

Mg+2

Na+1 Ne F-1 O-2 N-3

Bonding

What is a Model?

• Explains how nature operates.

• Derived from observations.

• It simplifies them and categorizes the information.

• A model must be sensible, but it has limitations.

Properties of a Model• A human inventions, not a blown up picture of

nature.

• Models can be wrong, because they are based on speculations and oversimplification.

• Become more complicated with age.

• You must understand the assumptions in the model, and look for weaknesses.

• We learn more when the model is wrong than when it is right.Find the energy for this

Localized Electron Model• Simple model, easily applied.

• A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.

Three Parts

1) Valence electrons using Lewis structures

2) Prediction of geometry using VSEPR

3) Description of the types of orbitals

VESPRNotes

Resonance Structures

• We have assumed up to this point that there is one correct Lewis structure.

• There are systems for which more than one Lewis structure is possible:• Different atomic linkages: Structural Isomers• Same atomic linkages, different bonding: Resonance

Resonance Structures (cont.)• The classic example: O3.

O O O

O O O

O O O

Both structures are correct!

Resonance Structures (cont.)• In this example, O3 has two resonance

structures:

O O O

• Conceptually, we think of the bonding being an average of these two structures.

• Electrons are delocalized between the oxygens such that on average the bond strength is equivalent to 1.5 O-O bonds.

Structural Isomers• What if different sets of atomic linkages can

be used to construct correct LDSs:

OCl Cl ClCl O

OCl Cl ClCl O

• Both are correct, but which is “more” correct?

Formal Charge• Formal Charge: Compare the nuclear charge (+Z) to the

number of electrons (dividing bonding electron pairs by 2). Difference is known as the “formal charge”.

OCl Cl ClCl O

#e- 7 6 7 7 6 7Z+ 7 6 7 7 7 6

Formal C. 0 0 0 0 +1 -1

• Structure with less F. C. is more correct.

Formal Charge• Example: CO2

OCO O O C O O C

e- 6 4 6 6 4 6 7 4 5

Z+ 6 4 6 6 6 4 6 6 4

FC 0 0 0 0 +2 -2 -1 +2 -1

More Correct

Beyond the Octet Rule

• There are numerous exceptions to the octet rule.

• We’ll deal with three classes of violation here:

• Sub-octet systems• Valence shell expansion• Odd-electron systems

Beyond the Octet Rule (cont.)• Some atoms (Be and B in particular) undergo bonding, but will form

stable molecules that do not fulfill the octet rule.

FBF

F

FBF

F

• Experiments demonstrate that the B-F bond strength is consistent with single bonds only.

Beyond the Octet Rule (cont.)• For third-row elements (“Period 3”), the energetic proximity of the d

orbitals allows for the participation of these orbitals in bonding.

• When this occurs, more than 8 electrons can surround a third-row element.

• Example: ClF3 (a 28 e- system)

FClF

F F obey octet rule

Cl has 10e-

Beyond the Octet Rule (cont.)• Finally, one can encounter odd electron systems where full pairs

will not exist.

ClO O

• Example: Chlorine Dioxide.

Unpaired electron

Summary• Remember the following:

• C, N, O, and F almost always obey the octet rule.• B and Be are often sub-octet• Second row (Period 2) elements never exceed the octet rule• Third Row elements and beyond can use valence shell expansion to exceed the octet rule.

• In the end, you have to practice…..a lot!

VESPR Background

• Recall from last lecture that we had two types of electron pairs: bonding and lone.

• The Lewis Dot Structure approach provided some insight into molecular structure in terms of bonding, but what about geometry?

• Valence Electron Shell Pair Repulsion (VESPR).

3D structure is determined by minimizing repulsion of electron pairs.

VESPR Background (cont.)

• Example: CH4

• Must consider both bonding and lone pairs in minimizing repulsion.

H C

H

H

H

Lewis Structure VESPR Structure

VESPR Background (cont.)

• Example: NH3 (both bonding and lone pairs).

Lewis Structure VESPR Structure

H N

H

H

VESPR Applications

• The previous examples illustrate the strategy for applying VESPR to predict molecular structure:

1. Construct the Lewis Dot Structure2. Arranging bonding/lone electron pairs

in space such that repulsions are minimized.

VESPR Applications

• Linear Structures: angle between bonds is 180°

F Be F

F Be F

• Example: BeF2

180°

VESPR Applications• Trigonal Planar Structures: angle between bonds is 120°

• Example: BF3

FBF

F

FBF

F

120°

VESPR Background (cont.)• Pyramidal: Bond angles are <120°, and structure is nonplanar:

H N

H

H

• Example: NH3

107°

VESPR Applications• Tetrahedral: angle between bonds is ~109.5°

• Example: CH4

H C

H

H

H

109.5°

VESPR Applications• Tetrahedral: angle may vary from 109.5° exactly due to size differences between bonding and lone pair electron densities

bonding pair

lone pair

VESPR Applications• Classic example of tetrahedral angle shift from

109.5° is water:

VESPR Applications• Comparison of CH4, NH3, and H2O:

VESPR Applications• Trigonal Bipyramidal, 120° in plane, and two

orbitals at 90° to plane:

P

Cl

Cl

Cl

Cl

Cl

• Example, PCl5:

90°

120°

VESPR Applications• Octahedral: all angles are 90°:

• Example, PCl6:

90°P

ClCl

Cl

Cl

ClCl

Advanced VESPR Applications

• Square Planar versus “See Saw”

See Saw

Square Planar

No dipole moment


• Driving force for last structure was to maximize the angular separation of the lone pairs.


• VESPR and resonance structures. Must look at VESPR structures for all resonance species to predict molecular properties.

O O O

O O OO

VESPR Applications

• Provide the Lewis dot and VESPR structures for CF2Cl2. Does it have a dipole moment?

C

F

FCl

Cl

32 e-

F

F

ClCl

Tetrahedral

Examples

• XeO3

• NO43-

• SO2Cl2

A useful equation• (happy-have) / 2 = bonds

• CO2 C is central atom

• POCl3 P is central atom

• SO42-

S is central atom

• SO32-

S is central atom

• PO43-

P is central atom

• SCl2 S is central atom

Exceptions to the octet• BH3

• Be and B often do not achieve octet• Have less than an octet, for electron

deficient molecules.

• SF6

• Third row and larger elements can exceed the octet

• Use 3d orbitals?

• I3-

Exceptions to the octet

• When we must exceed the octet, extra electrons go on central atom.

• (Happy – have)/2 won’t work

• ClF3

• XeO3

• ICl4-

• BeCl2

Formal Charge• For molecules and polyatomic ions that

exceed the octet there are several different structures.

• Use charges on atoms to help decide which.

• Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work.

Formal Charge• The difference between the number of

valence electrons on the free atom and that assigned in the molecule or ion.

• We count half the electrons in each bond as “belonging” to the atom.

• SO4-2

• Molecules try to achieve as low a formal charge as possible.

• Negative formal charges should be on electronegative elements.

Examples

• SiF4

• SeF4

• KrF4

• BF3

• PF3

• BrF3

No central atom

• Can predict the geometry of each angle.

• build it piece by piece.

How well does it work?

• Does an outstanding job for such a simple model.

• Predictions are almost always accurate.

• Like all simple models, it has exceptions.

• Doesn’t deal with odd electrons

Bonding

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